Part 1 The
kinetic
particle model and describing and explaining the properties of gases, liquids and solids, state changes
and solutions

EXAMPLES OF THE THREE PHYSICAL STATES OF
MATTER

GASES eg the air
mixture around us (including the oxygen needed for combustion) and the high
pressure steam in the boiler and cylinders of the steam locomotive. All of
these gases are 'invisible', being colourless and transparent, so note that
the 'steam' you see outside of the locomotive is actually fine liquid
droplets of water, formed from the expelled steam gas condensing when it
meets the cold air - the 'state change' of gas to liquid (same effect in
mist and fog formation).

LIQUIDS eg water is the
most common example, but so are, milk, hot butter, petrol, oil, mercury or
alcohol in a thermometer.

SOLIDS eg stone, all metals
at room temperature (except mercury), rubber of walking boots and the majority of physical
objects around you. In fact most objects are useless unless they have a
solid structure!

On this page the basic physical properties of gases,
liquids and solids are described in terms of structure, particle movement
(kinetic particle theory),
effects of temperature and pressure changes, and particle models used to explain these
properties and characteristics. Hopefully, theory and fact will match up to give students a clear
understanding of the material world around them in terms of gases, liquids and
solids - referred to as the three physical states of matter. The changes of
state known as melting, fusing, boiling, evaporating, condensing, liquefying,
freezing, solidifying, crystallising are described and explained with particle
model pictures to help understanding. There is also a mention of miscible and
immiscible liquids and explaining the terms volatile and volatility when applied
to a liquid.

When a gas is confined in a container the particles will cause and exert a
gas pressure which is measured in atmospheres (atm) or Pascals (1.0
Pa = 1.0 N/m2)
- pressure is force/area on which force is exerted.

The gas pressure is caused by the force created by millions of impacts of
the tiny individual gas particles on the sides of a container.

For example - if the number of gaseous particles in a container is doubled, the gas
pressure is doubled because doubling the number of molecules doubles the
number of impacts on the side of the container so the total impact force per
unit area is also doubled.

This doubling of the particle impacts doubling the pressure is pictured in
the two diagrams below.

If the volume of a sealed container is kept constant and the gas inside
is heated to a higher temperature, the gas pressure increases.

The reason for this is that as the particles are heated they gain kinetic
energy and on average move faster.

Therefore they will collide with the sides of the container with a
greater force of impact, so increasing the pressure.

There is also a greater frequency of collision with the sides of the
container BUT this is a minor factor compared to the effect of increased kinetic
energy and the increase in the average force of impact.

Therefore a fixed amount of gas in a sealed container of constant volume,
the higher the temperature the greater the pressure and the lower the
temperature the lesser the pressure.

If the ‘container’ volume can change, gases readily expand* on heating because of the lack of particle attraction, and
readily contract on cooling.

On heating, gas particles gain kinetic energy,
move faster and hit the sides of the container more frequently, and
significantly, they hit with a greater force.

Depending on the container situation, either or both of the pressure or volume will increase (reverse on cooling).

Note: *
It is the gas volume that expands NOT the molecules, they stay the same
size!

If there is no volume restriction
the expansion on heating is much greater for gases than liquids or solids
because there is no significant attraction between gaseous particles. The
increased average kinetic energy will make the gas pressure rise and so
the gas will try to expand in volume if allowed to e.g. balloons in a warm
room are significantly bigger than the same balloon in a cold room!

The natural rapid and random movement of the particles in
all directions means that gases readily ‘spread’ or
diffuse.

The net movement of a particular gas will be in the direction
from lower concentration to a higher concentration, down the so-called diffusion gradient.

Di

ffusion
continues until the concentrations are uniform throughout the container of
gases, but ALL the particles keep moving with their ever present kinetic energy!

Diffusion is faster in gases than
liquids where there is more space for them to move
(experiment illustrated below) and
diffusion is negligible in solids due to the close packing of the particles.

Diffusion is responsible for the
spread of odours even without any air disturbance e.g. use of perfume,
opening a jar of coffee or the smell of petrol around a garage.

The rate of diffusion increases with increase in temperature as the particles
gain kinetic energy and move faster.

Other evidence for random particle
movement including diffusion:

When smoke particles are viewed under a
microscope they appear to 'dance around' when illuminated with a light
beam at 90o to the viewing direction. This is because the
smoke particles show up by reflected light and 'dance' due to the
millions of random hits from the fast moving air molecules. This is
called 'Brownian motion' (see
below in liquids). At any given instant of time,
the hits will not be even, so the smoke particle get a greater bashing
in a random direction.

A two gaseous molecule
diffusion experiment is illustrated above and explained below!

A long glass tube (2-4 cm diameter) is filled at one
end with a plug of cotton wool soaked in conc. hydrochloric acid
sealed in with a rubber bung (for health and safety!) and the tube is
kept perfectly still, clamped in a horizontal position. A similar plug of
conc. ammonia solution
is placed at the other end. The soaked cotton wool plugs will give off
fumes of HCl
and NH3 respectively,
and if the tube is left
undisturbed and horizontal, despite the lack of tube movement, e.g. NO
shaking to mix and the absence of convection, a white cloud forms about
1/3rd
along from the conc. hydrochloric acid tube end.

Explanation: What happens is the colourless
gases, ammonia and hydrogen chloride, diffuse down the tube and
react to form fine white crystals of the salt ammonium chloride.

ammonia
+ hydrogen chloride
==> ammonium
chloride

NH3(g) + HCl(g)==> NH4Cl(s)

Note the rule: The smaller the
molecular mass, the greater the average speed of the molecules
(but all gases have the same average kinetic energy at the same
temperature).

Therefore the smaller the
molecular mass, the faster the gas diffuses.

A
coloured
gas, heavier than air (greater density), is put into the
bottom gas jar and a second gas jar of lower density colourless air is placed over it separated with a
glass cover.

If the glass cover is removed
then (i) the colourless air gases diffuses down into the coloured brown
gas and (ii) bromine diffuses up into the air. The particle movement
leading to mixing cannot be due to convection because the more dense gas starts at the
bottom!

No 'shaking' or other means of mixing is required. The
random movement of both lots of particles is enough to ensure that both gases
eventually become completely mixed by diffusion.

This is clear evidence for diffusion
due to the random continuous movement of all the gas particles and,
initially, the net movement of one type of particle from a higher to a
lower concentration ('down a diffusion gradient'). When fully mixed, no further colour change
distribution is observed BUT the random particle movement continues! See
also other evidence in the liquid section below.

A note on 'forces'

Forces between particles are
mentioned on this page and some ideas will seem more abstract than others - but think
about it ...

A gas spreads everywhere
in a given space, so there can't be much attraction between the
molecules/particles.

Something must hold liquid
molecules together or how can a liquid form from a gas?

In fact between liquid molecules there are actually weak electrical forces of attraction
called intermolecular forces, but they can't be strong enough to create a rigid solid
structure.

However, in solids, these forces must be
stronger to create the rigid structure.

How does the
kinetic particle theory of liquids explain the properties of liquids?

A liquid has a fixed volume at a given temperature but
its shape is that of the container which holds the liquid.

There are m

uch greater forces of attraction between the particles in a liquid compared to
gases, but not quite as much as in solids.
If there were no intermolecular forces, liquids could not exist.

The particles are quite close together but still arranged at random throughout the container
due to their random movement, there is a little close range order as you can get clumps of particles clinging together temporarily
(as in the diagram above).

The p

articles are moving rapidly in all directions but
collide more frequently with each other than in gases due to
shorter distances between particles.

With increase in temperature, the particles
move faster as they gain kinetic energy, so increased
collision rates, increased collision energy, increased rates of particle diffusion,
expansion leading to decrease in density.

Liquids have a much greater density than gases (‘heavier’) because the particles are much closer
together because of the attractive forces.

Most liquids are just a little less dense than when they are
solid

Water is a curious exception to this general rule, which is
why ice floats on water.

Liquids usually flow freely despite the forces of attraction between the particles but liquids are not as ‘fluid’ as gases.

Note 'sticky' or viscous liquids have much stronger attractive forces between
the molecules BUT not strong enough to form a solid.

Liquids have a surface, and a fixed volume (at a particular temperature) because of the increased particle attraction, but the shape is not fixed and is merely that of the container itself.

Liquids seem to have a very weak 'skin' surface effect which is caused by
the bulk molecules attracting the surface molecules disproportionately.

Liquids are not readily compressed because
there is so little ‘empty’ space between the particles, so increase in pressure
has only a tiny effect on the volume of a solid, and you need a huge increase in
pressure to see any real contraction in the volume of a liquid.

Liquids will expand on heating but nothing like as much as gases because of the greater particle attraction restricting the
expansion (will contract on cooling).

Note: When heated, the liquid particles gain kinetic energy and hit the sides of the container more frequently, and more significantly, they hit with a greater force, so in a sealed container the pressure produced can be considerable!

The natural rapid and random movement of the particles means that liquids ‘spread’ or
diffuse. Diffusion is much slower in liquids compared to gases because there is less space for
the particles to move in and more ‘blocking’ collisions happen.

Just dropping lumps/granules/powder of a soluble solid (preferably
coloured!) will resulting in a dissolving followed by an observable diffusion
effect.

Again, the net flow of dissolved particles will be from a higher
concentration to a lower concentration until the concentration is uniform
throughout the container.

If coloured crystals of e.g. the highly
coloured salt crystals of potassium manganate(VII) are dropped into a
beaker of water and covered at room temperature.

Despite the lack of
mixing due to shaking or convection currents from a heat source etc. the bright purple colour of the dissolving salt
slowly spreads throughout all of the liquid but it is much slower than the
gas experiment described above because
of the much greater density of particles slowing the spreading due to close
proximity collisions.

The same thing happens with dropping
copper sulphate crystals (blue, so observable) or coffee granules into water and just leaving the
mixture to stand.

When pollen grains are viewed under a
microscope they appear to 'dance around' when illuminated with a light
beam at 90o to the viewing direction.

This is because the
pollen grains show up by reflected light and 'dance' due to the
millions of random hits from the fast moving water molecules.

This phenomenon is
called 'Brownian motion' after a
botanist called Brown first described the effect (see
gases above).

At any given instant of time,
the hits will not be even all round the pollen grain, so they get a greater
number of hits in a random direction.

We can use the state particle models and diagrams to explain changes of state and the energy changes involved.

These are NOT chemical changes BUT PHYSICAL CHANGES, e.g.
the water molecules H2O are just the same in ice, liquid water, steam
or water vapour. What is different, is how they are arranged, and how strongly
they are held together by intermolecular forces in the solid, liquid and gaseous
states.

Because of random collisions, the particles in
a liquid have a variety of speeds and kinetic energies. On heating, particles gain kinetic energy
and move faster and are more able to overcome the intermolecular forces
between the molecules i.e. some particles will have enough kinetic energy to
overcome the attractive forces holding the particles together in the bulk
liquid.

Even without further heating, evaporation occurs all the time from
volatile liquids, but it is still the higher kinetic energy particles that can
overcome the attractive forces between the molecules in the bulk of the
liquid and escape from the surface into the surrounding air.

In evaporation* and boiling
(both are vaporisation) it is the
highest kinetic energy molecules that can ‘escape’ from the attractive forces of the other liquid particles.

The particles lose any order and
become completely
free to form a gas or vapour.

Energy is needed to overcome the attractive forces between particles in the liquid and is taken in from the surroundings.

This means heat is taken in,
so evaporation and boiling are endothermic processes
(ΔH +ve).

If the temperature is high enough boiling
takes place.

Boiling is rapid evaporation
anywhere in the bulk liquid and at a fixed temperature called the boiling point and requires continuous addition of heat.

The rate of boiling is limited by the rate
of heat transfer into the liquid.

* Evaporation takes place more slowly
than boiling
at any temperature between the melting point and boiling point,
and only from the
surface, and results in the liquid becoming cooler due to
loss of higher kinetic energy particles.

More details on the e

nergy changes for these physical changes of state
for a range of substances are dealt with in a section of
theEnergetics Notes.

On cooling, liquid particles lose kinetic energy and so can
become more strongly attracted to each other.

When the temperature is low enough, the kinetic energy of the particles is
insufficient to prevent the particle attractive forces causing a solid to
form.

Eventually at the freezing point the forces of attraction are sufficient to remove any remaining freedom
of movement (in terms of one place to another) and the particles come together to form the ordered solid arrangement
(though the particles still have vibrational kinetic energy.

Since heat must be removed to the surroundings,
so strange as it may seem,
freezing is an exothermic process (ΔH -ve).

and the comparative energy changes of state changes gas <=>
liquid <=> solid

2f(i)
Cooling curve:
Note the temperature stays constant during the state changes of condensing
at temperature Tc, and freezing/solidifying at temperature Tf.
This is because all the heat energy removed on cooling at these temperatures
(the
latent heats or enthalpies of state change), allows
the strengthening of the inter-particle forces without temperature fall (the
heat loss is compensated by the exothermic increased intermolecular
force attraction).
In between the 'horizontal' state change sections of the graph, you can
see the energy 'removal' reduces the kinetic energy of the particles,
lowering the temperature of the substance.

Energy changes
for these physical changes of state for a range of substances are dealt with in a section of
the Energetics Notes.

2f(ii)Heating curve:
Note the temperature stays constant during the state changes of melting
at temperature Tm and boiling at temperature Tb. This is because all the energy absorbed in
heating at these temperatures
(the latent heats or enthalpies of state change),
goes into weakening the inter-particle
forces without temperature rise
(the heat gain equals the endothermic/heat absorbed energy required to
reduce the intermolecular forces). In between the 'horizontal' state change
sections of the graph, you can see the energy input increases the
kinetic energy of the particles and raising the temperature of the
substance.

This is when a
solid, on heating, directly changes into a gas without melting, AND the
gas on cooling re-forms a solid directly without condensing to a
liquid. They usually involve
just a physical change BUT its not always that simple!

Theory in terms of particles:

When the solid is heated
the particles vibrate with increasing force from the added thermal
energy.

If the particles have enough kinetic energy of vibration to
partially overcome the particle-particle attractive forces you would
expect the solid to melt.

HOWEVER, if the particles at this point have
enough energy at this point that would have led to boiling, the liquid
will NOT form and the solid turns directly into a gas.

Eventually, when the particle kinetic energy is low
enough, it will allow the particle-particle attractive forces to produce
a liquid.

BUT the energy may be low enough to permit direct formation of
the solid, i.e. the particles do NOT have enough kinetic energy to
maintain a liquid state!

Overall exothermic
change, energy released and 'given out' to the surroundings.

Examples:

Even at room temperature
bottles of solid iodine show crystals forming at the top of the bottle
above the solid. The warmer the laboratory, the more crystals form when
it cools down at night!

I2 (s)
I2 (g) (physical
change only)

The formation of a particular
form of
frost involves the direct freezing of water vapour (gas).
Frost can also evaporate directly back to water vapour (gas) and this
happens in the 'dry' and extremely cold winters of the Gobi Desert on a
sunny day.

H2O
(s)
H2O (g) (physical change only)

Solid carbon dioxide (dry ice)

is
formed on cooling the gas down to less than -78oC. On warming
it changes directly to a very cold gas!, condensing any water vapour in
the air to a 'mist', hence its use in stage effects.

CO2
(s)
CO2 (g) (physical change only)

On heating strongly in a test tube, the
white solid ammonium chloride, decomposes into a mixture of two
colourless gases ammonia and hydrogen chloride. On
cooling the reaction is reversed and solid ammonium chloride reforms at
the cooler top of the test tube.

Ammonium chloride

+
heat energy ammonia + hydrogen chloride

NH4Cl(s)
NH3(g) + HCl(g)

This involves both chemical and
physical changes and is so is more complicated than examples 1. to 3. In
fact the ionic ammonium chloride crystals change into covalent
ammonia and hydrogen chloride gases which are naturally far more
volatile (covalent substances generally have much lower melting and
boiling points than ionic substances).

PLEASE NOTE, At a higher
level of study (e.g. UK A2 advanced level),
you need to study
the g-l-s phase diagram for water and the vapour pressure curve of ice at
particular temperatures. For example, if the ambient vapour
pressure is less than the equilibrium vapour pressure at the temperature of
the ice, sublimation can readily take place. The snow and ice in the Gobi
Desert do not melt in the Sun, they just slowly 'sublimely' disappear!

3a.WHAT HAPPENS TO PARTICLES WHEN A
SOLID DISSOLVES IN A LIQUID SOLVENT?

What do the words
SOLVENT, SOLUTE and SOLUTION mean?

When a solid (the
solute) dissolves in a liquid (the
solvent) the resulting
mixture is
called a solution.

In general: solute +
solvent ==> solution

So, the solute is what dissolves in a solvent, a solvent is
a liquid that dissolves things and the solution is the result of dissolving
something in a solvent.

The solid loses all its
regular structure and the individual solid particles (molecules or ions) are
now completely free from each other and randomly mix with the
original liquid particles, and all particles can move around at random.

This describes salt
dissolving in water, sugar dissolving in tea or wax dissolving in a
hydrocarbon solvent like white spirit.

It does not usually
involve a chemical reaction, so it is generally an example of a physical
change.

Whatever the changes in
volume of the solid + liquid, compared to the final solution, the Law of
Conservation of Mass still applies.

This means: mass of
solid solute + mass of liquid solvent = mass of solution after mixing
and dissolving.

You cannot create mass
or lose mass, but just change the mass of substances into another form.

If the solvent is
evaporated, then the solid is reformed e.g. if a salt solution is left
out for a long time or gently heated to speed things up, eventually salt
crystals form, the process is called crystallisation.

3b.WHAT HAPPENS TO PARTICLES WHEN
TWO LIQUIDS COMPLETELY MIX WITH EACH OTHER?

WHAT DOES THE WORD
MISCIBLE MEAN?

If two liquids completely
mix in terms of their particles, they are called miscible liquids
because they fully dissolve in each other. This is shown in the diagram below
where the particles completely mix and move at random. The process can be
reversed by
fractional distillation.

+

3c.WHAT HAPPENS TO PARTICLES WHEN
TWO LIQUIDS DO NOT MIX WITH EACH OTHER?

WHAT DOES THE WORD
IMMISCIBLE MEAN?

WHY DO THE LIQUIDS
NOT MIX?

If the two liquids do NOT
mix, they form two separate layers and are known as immiscible liquids,
illustrated in the diagram below where the lower purple liquid will be more
dense than the upper layer of the green liquid.

You can separate these two
liquids using a separating funnel.

The reason for this is that the
interaction between the molecules of one of the liquids alone is stronger
than the interaction between the two different molecules of the different
liquids.

For example, the force of attraction between water molecules is
much greater than either oil-oil molecules or oil-water molecules, so two
separate layers form because the water molecules, in terms of energy change,
are favoured by 'sticking together'.